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Equilibrium: A state of dynamic balance

Equilibrium: A state of dynamic balance. I will recognize the characteristics of chemical equilibrium. I will write equilibrium expressions for systems that are at equilibrium. I will calculate equilibrium constants from concentration data. completion. Reaction goes to completion

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Equilibrium: A state of dynamic balance

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  1. Equilibrium: A state of dynamic balance • I will recognize the characteristics of chemical equilibrium. • I will write equilibrium expressions for systems that are at equilibrium. • I will calculate equilibrium constants from concentration data.

  2. completion • Reaction goes to completion • When a reaction results in almost complete conversion of reactants to products • Rarely happens

  3. Reversible reaction • Most reactions • Do NOT go to completion • Appear to stop • Are reversible • Reversible reaction • One that can occur in both the forward and the reverse directions • Denoted with a double arrow to show that both reactions occur • Forward reaction = reactants on left • Reverse reaction = reactants on right

  4. Chemical equilibrium Chapter 18 I will discover that many reactions and processes reach a state of chemical equilibrium. I will use Le Chatelier’s Principle to explain how various factors affect chemical equilibrium. I will calculate equilibrium concentrations of reactants and products using the equilibrium constant expression. I will determine the solubilities of sparingly soluble ionic compounds.

  5. Reversible Reactions

  6. Rate of reactionDepends on concentration of the reactants N2 (g) + 3 H2 (g) < -- > 2NH3 (g) • The concentrations of reactants decrease at first. • The concentrations of the product increases at first. • Then, before all the reactants are used up, all concentrations become constant.

  7. Chemical equilibrium • A state in which the forward and reverse reactions balance each other because they take place at equal rates • Rate forward reaction = Rate reverse reaction • Concentrations of reactants and products are constant • HOWEVER! • The amounts or concentrations of reactants and products • Are NOT usually equal • MAY even differ by a factor of a million or more!

  8. Equilibrium expressions and constants • Majority of reactions reach equilibrium with varying amounts of reactants unconsumed • NOT all our predicted moles of product gets produced • Law of Chemical Equilibrium • At a given temperature, a chemical system may reach a state in which a particular ratio of a reactant and product concentrations has a constant value • aA + bB<--> cC + dD

  9. Equilibrium Constant • The numerical value of the ratio of product concentrations to reactant concentrations • Constant only at a SPECIFIC TEMPERATURE • Products on top, reactants on bottom • Keq > 1: MORE products than reactants at equilibrium • Keq < 1: LESS products than reactants at equilibrium • Keq = [C]c[D]d [A]a[B]b

  10. HOMOgeneousequilibrum • All the products and reactants are in the same physical state • Must use ALL CONCENTRATIONS for Keq

  11. Heterogeneous Equilibrium • Reactants and products of a reaction are present in more than one physical state • Do NOT count concentrations of solids or liquids when calculating Keq • Can be OMITTED from the Keq expressions • If a solid or liquid state of a substance is present in addition to the gas state….LABEL the gas concentration in your expression to distinguish between the two

  12. Factors affecting chemical equilibrium • I will describe how various factors affect chemical equilibrium. • I will explain how Le Chatelier’s Principle applies to equilibrium systems.

  13. Le Chatelier’s Principle • Apply stress to a system at equilibrium • System will shift in the direction that relieves the stress • Stress • Any kind of change in a system at equilibrium that UPSETS equilibrium • Types: • Change in concentration • Change in volume (pressure) • Change in temperature

  14. Change in Concentration • Changes equilibrium POSITION • Shifts left or right • Does NOT change equilibrium constant (Keq) • Add reactant = shift right • Remove reactant = shift left • Add product = shift left • Remove product = shift right

  15. Change in Volume (Pressure) • Changes equilibrium POSITION • Shifts left or right • ONLY if # moles of gaseous reactants is DIFFERENT than # moles gaseous products • Does NOT change equilibrium constant (Keq) • Volume Pressure

  16. Change in Volume (Pressure) Decrease volume (increase pressure) • Situation 1: more moles gas reactants & less moles gas products • Shift right • Situation 2: moles gas reactants = moles gas products • NO shift • Situation 3: less moles gas reactants & more moles gas products • Shift left

  17. Change in Temperature • Changes equilibrium POSITION • Shifts left or right • CHANGES equilibrium constant (Keq) • Large Keq = more product in equilibrium mixture • Small Keq = less product in equilibrium mixture

  18. Change in Temperature Endothermic- absorbs heat Reactants + heat <--> Products Exothermic- releases heat Reactants <--> Products + heat • Hot • -∆H° (lose heat) • Forward reaction = exo, backward = endo • ↑temp = shift left, ↓Keq • ↓temp = shift right, ↑Keq • Cold • +∆H° (gain heat) • Forward reaction = endo, backward = exo • ↑temp = shift right, ↓Keq • ↓temp = shift left, ↑Keq

  19. Catalysts • Speeds up a reaction • Speeds it up EQUALLY in BOTH directions (Right & Left) • Helps a reaction reach equilibrium quickly • But NO CHANGE in the AMOUNT of PRODUCT formed

  20. Using equilibrium constants • I will determine equilibrium concentrations of reactants and products. • I will calculate the solubility of a compound from its solubility product constant. • I will explain the common ion effect.

  21. Using Equilibrium constants • Review: • Large Keq = Products favored • Small Keq = Reactants favored • Knowing the size of the Keq helps a chemist • Decide whether a reaction is practical for making a particular product • Calculate the equilibrium concentration of ANY substance involved in the reaction • (if the concentrations of all other reactants/products are known)

  22. Calculating Equilibrium Concentrations • Write the equilibrium constant (Keq) expression • Solve the equation for the unknown (using algebra skills) • Substitute in all known concentrations and the Keq value • Use calculator to find unknown concentration • Chemists would then use this concentration to determine if enough of their desired unknown could be produced in the reaction

  23. Calculating Equilibrium Concentrations

  24. Solubility Equilibria Ionic Compounds Some dissolve readily in water • Ex NaCl(s) • High solubility Some barely dissolve at all • Ex BaSO4(s) • low solubility

  25. Solubility Product Constant • An equilibrium constant for the dissolving of a sparingly soluble ionic compound in water • Ksp = the product of the concentrations of the ions each raised to the power equal to the coefficient of the ion in the chemical equation • Small Ksp = Products NOT favored at equilibrium

  26. Solubility Product Constant • Example: (Remember it depends ONLY on [ IONs])

  27. Review: Solubility • Solubility in water • The amount of the substance (moles) that will dissolve in a given volume of water (Liter)

  28. Molar Solubility/calculating Ion Concentration • To determine solubility of a sparingly soluble compound X X 2X S = = (x)(2x)2

  29. Molar Solubility/ Calculating ion concentration from Ksp • 2) Use Ksp to calculate the following ion concentrations. • [F-] in a saturated solution of CaF2 • [Ag+] in a solution of AgBr at equilibrium • [Ag+] in a solution o f Ag2CrO4 at equilibrium 1) Calculate the solubility (Mol/L) of the ionic compound (@298K)

  30. Predicting Precipitates • Ksp can predict if a precipitate will form when two ionic compounds are mixed • Precipitate likely to form ONLY if either product has LOW solubility (small Ksp) • If concentrations of the ions (right side) are greater than the concentrations of the ionic compound (left side) then the reaction will precipitate (shift to the left) • JUST look at your pink sheet! • Insoluble = low solubility = small Ksp = precipitate will form!

  31. Common Ion effect • Common Ion • Ion common to two or more ionic compounds • Common Ion Effect • The lowering of the solubility of a substance by the presence of a common ion

  32. Common Ion Effect Left --common SO42-/Ba2+ lowers BaSO4 Solubility • Common ion effect will cause a shift in the equilibrium. • BaSO4(s) <-->Ba2+(aq) + SO42- (aq) • If you add Na2SO4 which direction will the equilibrium shift? •  If you add BaCl2 which direction will the equilibrium shift? • If you add NaF which direction will the equilibrium shift? • If you add NaCl which direction will the equilibrium shift? NaSO4 (s) BaF2 (I) Ba 2+reacts with SO42- and F-. So Ba used up twice as fast (need more…right) NaSO4 (s) BaCl2 (s) NO SHIFT

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